Jan.,
1921
THE JOURNAL OF I N D U S T R I A L A N D ENGIXEERING CHEMISTRY
perience with vitamines is of a similar character. It is, therefore, apparent that the two groups of substances conduct themselves, in respect t o fractionation procedures, in an entirely analogous manner. The one outstanding characteristic of an enzyme, which should serve to difr‘erentiate it from everything with which i t might be confused, is the property of accelerating chemical reactions, without itself being destroyed. This has been demonstrated, to a certain extent a t least, in the case of some of the wellcharacterized enzymes. That i t can be shown in the case of vitamines, however, is out of the question at present, since the only test of the activity of a vitamine is by means of a living organism, and in such cases the recovery of the vitamine a t the conclusion of its period of action is obviously impossible. There is this in common, however, that the apparent amount of vitamine required for a given result is of the same order of magnitude as required for the transformations effected by enzymes. Thus, for instance, i t has been shown that invertase can hydrolyze zoo,ooo times its weight of saccharose, and rennet can clot 400,ooo times its weight of caseinogen in milk. I n the case of vitamine, fractions have been prepared of which only a few tenths of a milligram per day are sufficient to supply the requirements of a pigeon maintained on a vitamine-free diet. On the basis of the above comparison i t is seen that, aside from a possibly significant degree of dialyzability, there is no outstanding evidence that vitamines should not be classed with the enzymes. This viewpoint is further strengthened by the negative evidence that, even in spite of the repeated efforts of able investigators, the original conception that vitamine is a well-characterked chemical individual capable of being isolated has never been realized. In conclusion, therefore, the question may well be raised as to whether our knowledge of vitamines will not be more rapidly advanced by tentatively including them in the class of substances designated as enzymes. THE MECHANISM OF CATALYTIC PROCESSES‘ By Hugh S. Taylor PRINCETONUNIVERSITY,PRINCETON,NEW JERSEY
HETEROGENEOUS CATALYSIS In reviewing the general field of contact catalysis, attention cannot but be directed to the diversity of views obtaining in reference to the mechanism of the process, many of which are capable of direct experimental check, which, unfortunately, in so many cases, is not applied. Sabatierl suggests t h a t hydrogenation and dehydrogenation processes occurring in contact with finely divided metals are to be ascribed t o the capacity of these metals t o form unstable hydrides which interact with the other components of the system t o yield the reaction products. Thus, for the catalytic hydrogenation of ethylene in coatart with nickel, Sabatier suggests the following scheme: H1 Ni2 = NipHe NizHr C2H4 = CPHB Nil Baiwrofts suggests that it seems natural t o assume that the selective adsorption of the reaction products is the determining factor. This conclusion, however, Bancroft shows, is not entirely satisfactory in view of the known experimental behavior of certain reactions studied. Thus, ethylene can be produced by catalytic dehydration of alcohol by means of alumina even in the presence of a large amount of water vapor. The beautiful studies of catalytic actions a t solid surfaces recently made by Armstrong and Hilditch4 lead to a conclusion which is the
+
+
+
1 Abstract of a lecture given before the New York Section of the American Chemical Society, December 10, 1920. “ L a Catalyse e n Chimie Organique,” 2nd Edition, 1920, p. 60. 3 Presidential Address, American Electrochemical Society, April 1920. PYOC. R o y . Soc., 96 (1919), 137, 322; 97 (1920), 259, 265; 98 (1920), 2 : .
75
antithesis of the views of Sabatier. Armstrong and Hilditch are inclined to regard the affinity of the carbon compound rather than that of the hydrogen to the metal as of prime importance, indeed, as the determining factor. In the hydrogenation of unsaturated oils their experimental data lead them to the conclusion that the process of catalytic hydrogenation in the solidliquid state involves the primary formation of an unstable complex or “intermediate compound” between nickel and the unsaturated compound. Dehydration reactions subsequently studied lead them to similar conclusions in reference to primary formation of nickel-organic compound complexes. Lewis1 assumes that the mechanism of hydrogenation involves essentially the dissociation of hydrogen, either adsorbed on or absorbed by the nickel, followed by collisions between the charged nickel particles and the unsaturated molecules. He concludes that, in the case of hydrogenation of olein and of similar substances, adsorption of the unsaturated compound on the metal does not take place, the adsorption being restricted t o metal hydrogen components. hlany observations made in the course of experimental work a t Princeton tend to show that. in the case of a variety of different substances there occurs a definitely measui able adsorption by catalytic agents of one or other of the reactants in a catalytic change. In the study of the reaction kinetics of various catalytic processes, indirect evidence has led to the conclusion that, iizter alia, benzene vapor is strongly adsorbed by nickel, and carbon monoxide by nickel at temperatures as high as 150’ C. Carbon dioxide is apparently adsorbed by iron oxide a t temperatures up t o 250’ C. Water vapor is adsorbed by various metal catalysts. Systematic study of the magnitude of the adsorption effect with a series of gases and a variety of catalytic agents has, therefore, been undertaken. The preliminary results obtained are remarkable and serve to show the advances in our knowledge of mechanism of catalytic change which may come from such experimental study. NICKEL-with Mr. A. W. Gauger, the adsorptions by nickel of hydrogen, carbon monoxide, carbon dioxide, and ethylene, using nitrogen as the reference gas have been determined in the temperature ranges in which these gases react with one another. The material used was reduced nickel on a porous support of Non-Pareil Diatomite Brick, 7.5 g. of the material being employed, containing 0.75 g. of metallic nicke!. The porous support used was graded between 8- and Io-mesh sieves. Table I shows the cubic centimeters of different gases measured at oo C. and 760 mm. pressure which were required to fill the vessel containing the nickel catalyst at 760 mm. pressure and various temperatures. TABLEI GAS
---Temperature 21 175 15.04 9.8 13.6 11.1 14.05 14.07
............ ................
Nitrogen.. Hydrogen Carbon dioxide Carbon monoxide Ethylene..
. ................ .. . ...............
of Absorption Vessel, C.200 225 250 275 9.4 8.8 8.5 8.1 13.2 ... 12.0 ... 10.6 9.9 9.4 ...
......
...... ......
......
If it be assumed that the adsorption of nitrogen by nickel is negligible, the following values for the adsorption of different gases, per gram of nickel upon the given porous support, are readily derived. ADSORPTION I N C C . (AT 0’ Temperature C. 175 Hs.. 5.2 COa 1.7 CO 5.66 CzHa. . . . . . . . . . . 6 . 5
c. AND
...........
........... ............
760
MM.)
200 5.1
225
1.6
.. .. ..
... ...
...
1.5
PER GRAM
Ni
250 4.73 1.33
...
...
With ethylene, only one set of experimental measurements has, as yet, been made. It suffices, however, t o show that this gas is more adsorbed than any of the other gases studied. With carbon monoxide, the measurements have been limited to the one temperature because, a t lower temperatures, the question 1
J. Chem. S O C , 117 (1920), 623.
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T H E JOURNAL OF I N D U S T R I A L A N D ENGINEERING CHEMISTRY
of the formation of nickel carbonyl, Ni(CO)r, would necessarily intrude. Above the temperature of 175' C. the measurements are complicated by the catalytic decomposition of carbon monoxide to form carbon and carbon dioxide
zco
=
coz + c.
The adsorption of carbon dioxide is noteworthy, although smaller in magnitude than that of the other gases studied. With hydrogen, the initial adsorption effect is followed by a secondary slow solubility effect which causes a slow increase in the volume of gas required to fill the reaction vessel. This secondary change is, however, so slow that the initial adsorption effect can readily be measured with an accuracy of I per cent. The results thus obtained with nickel may be generalized. The gases which take part in the following reactions: CO 3H2 = CHI H?O COz 4H2 = CHL aH,O CZH4 3- HZ = CZH6 are all markedly adsorbed by a nickel catalyst in t h t temperature range in which they react to form the stated reaction products. coPPsR-Similar experiments with these gases have been performed by Mr. R. M. Burns, employing copper obtained by reduction, at low temperatures, of copper oxide. The oxide was produced by calcination of the nitrate in a stream of air. The interest attaching to this study arises from the observations of Sabatier with respect to copper as a catalytic agent. Sabatier states that, under no conditions, can copper induce the interaction of carbon monoxide or carbon dioxide with hydrogen to form methane. On the contrary, above 160' C., ethylene and hydrogen react in contact with copper to yield the saturated hydrocarbon, ethane. Preliminary experiments showed that the adsorption effects with this metal were of a much lower order of magnitude than with nickel. Consequently a larger sample of reduced metal, 22.9 g., was used for the determinations. The measurements of adsorption were made at 25 O. C., I 10' C,, and 218' C., a t a pressure of 760 mm. The gases studied were again nitrogen, carbon monoxide, carbon dioxide, hydrogen, and ethylene. As a check on the nitrogen determination, to show that the figures obtained with this gas represented zero adsorption, one determination was made a t 25' C., with a specially purified sample of helium, obtained through the courtesy of the U. S. Bureau of Mines. Table I1 shows the number of cubic centimeters of the different gases (measured at oo C. and 760 mm. pressure) which are required t o fill the reaction vessel containing the reduced copper, when this is maintained a t the three stated temperatures
+ +
+ +
TABLEI1 Cc. G a s Required t o Fill Ve!jsel a t 25' C. 1100 c. 218' C. GAS Helium. 22.35 Nitrogen.. .............. 2 2 . 4 l?:46 1i:9 Hydrogeq.. ............. 2 2 . 4 17.6 13.9 Carbon dioxide.. . . . . . . . . . 22.55 17.5 13.9 Carbon monoxide.. ...... 2 3 . 9 18.1 13.9 Ethylene.. 24.1 18.1 13.9
................ ..............
The experiments show that only with ethylene and carbon monoxide is there a measurable adsorption and with these gases only a t the two lower temperatures. At the temperature of 218' C., the volume of gas adsorbed is immeasurably small in every case. The experiments with copper and with nickel both show, therefore, a greater adsorption of the unsaturated compound than of hydrogen. It is the view of Armstrong and Hilditch rather than that of Sabatier and Lewis which the present experimental observations, therefore, tend to support, though naturally a wide extension of the experimental range will be necessary before any definite conclusions can be reached. This extension is in progress. We are engaged on measurements of adsorption with a wide variety of metals and metallic oxides under varied conditions.
Vol. 13, No. I
In connection with the adsorption experiments with ethylene on copper, it is interesting to note that a t the temperature a t which hydrogenation commences (160') the adsorption of ethylene is already quite low. In other words, a t this temperature, the ethylene evaporates rapidly from the copper surface after condensation has occurred. The experimental results obtained with the gas a t lower temperatures show that the copper surface must be relatively free from adsorbed ethylene a t the temperature of hydrogenation. This is probably true also in the case of the nickel experiments previously described. This factor appears to us to be of cardinal importance in a discussion of the mechanism of contact action. Furthermore, the fact that, as far as adsorption by copper is concerned, carbon monoxide behaves like ethylene, whereas hydrogenation of carbon monoxide in contact with copper cannot be achieved, shows that further insight into the several factors prevailing is still needed. We propose t o obtain this by extending our studies on adsorption by various metallic catalysts which either promote or are inert in the hydrogenation process. Thus, in contact with cobalt, carbon monoxide and hydrogen yield methane. With iron, no methane is obtained.' Since carbon monoxide and hydrogen do not interact in coiltact with reduced copper it is possible to study the adsorption of these gases from mixtures of the same. Similar studies can be carried out with mixtures of ethylene and hydrogen a t temperatures below those a t which these gases interact. In a preliminary manner we have studied the adsorption of various mixtures of hydrogen and carbon monoxide and hydrogen and ethylene at 25' C. The results obtained are very remarkable and promise further insight into the catalytic process. In Table I11 are given the adsorptions in cubic centimeters of gas absorbed by 22.9 g. of reduced copper with various mixtures of the two pairs of gases. In the last column are given the calculated values for adsorption, if the amounts adsorbed were in direct proportion to the partial pressures of the gases present. TABLEI11 Cc. Gas Calculated Adsorption (at ' 0 760 Mm.) if Proportional
Absorb'ed a t 25' C. MIXTURE and 760 Mm. 0% Hz 100% CO 1.5 50% $2, 50% C O . . . . . . . . . . . 1 . 3 84.5% Hz. 15.5% C O . , . . . . . . 0 . 9 I 0 0 7 Hz ................... 0 . 0 0% I"rZ 100% CZH4.. . . . . . . . . 1 . 7 53% H z , 47% CzHi.. 1.2 100% Ha.. 0.0 GAS
...........
to Partial Pressures of Gases
........
.................
o:js 0.23
..
0 :8
It is thus apparent that carbon monoxide and ethylene are much more markedly adsorbed a t lower pressures than a t higher pressures, the adsorption tending to become independent of the pressure as this increases. THE ICINETICS OF CATALYTIC ACTIONS
I
The abnormal variation of adsorption with pressure constitutes a factor of considerable importance in regard to the mechanism of the catalytic process. If the catalytic reaction occurs in the surface layer it is apparent that the pressure-adsorption ratio determines the concentration of the reactants in the active layer. For example, in the reaction CzH4 Hz = C ~ H B the rate of formation of ethane in the gas phase is
+
Rt = ~ ( P c ~ H J ( P H J , where pea" and f i are ~ ~t h e partial pressures of the interacting gases. Similarly a t the surface of the copper, the rate of reaction is Ra = ~ ~ ( C C Z(CH2)r H~) where CC~H, and C H ~ are the concentrations of the gases a t the surface. Now, if the experimental conditions were so chosen that the concentration of ethylene in the surface layer was inde1
Sabatier, LOG.dt.
T H E J O U R N A L O F I N D U S T R I A L A N D EAVGIXEERILVGC H E M I S T R Y
Jan., 1 9 2 1
pendent of the prevailing partial pressure of the gas, i. e . ,
CC~H = ~~ ( P c ~ H ~=) O k, the reaction in the surface layer would become Rz = k z . k . ( C ~ ~ ) .
So, iP the hydrogen concentration were governed by Henry’s law, the reaction would be bimolecular in the gas phase and apparently monomolecular in the surface layer, The same considerations might be extended to the question of the equilibrium constant of the given reaction. In the case cited, with the same assumptions as to the distribution ratio between gas and surface layer, the equilibrium constant K, in the gas phase would be
In the surface layer, however, if the hydrogen and ethane obeyed Henry’s law, but the ethylene concentration was independent of the partial pressure of ethylene, the equilibrium constant, Ks+would be
77
quoted. Fink’s results on sulfur trioxide formation show no agreement with a termolecular reaction equation in the early stages of an experiment. Towards the completion of the process, however, an excellent termolecular constant, ks, is obtained as Table IV shows. On the interpretation given in the preceding paragraphs the distribution of sulfur dioxide and oxygen between the gas phase and the contact material must in the later stages of the reaction follow Henry’s law. Bodenstein and Ohlmer found that the reaction between oxygen and carbon monoxide in contact with quartz glass takes place a t a rate proportional to the pressure of oxygen and inversely proportional to the pressure of carbon monoxide. In contact with crystalline quartz, however, the reaction followed the ordinary stoichiometric equation, a result which should have attracted a much greater attention in the discussion of catalysis than it has yet done. On the interpretation here given, this diversity of reaction mechanism, in the same reaction, with the two catalysts, is to be ascribed to the different distribution ratios between the gas phase and the surface layer on the contact mass. An experimental test of such a viewpoint could be carried out.
It is apparent, therefore, that the position of equilibrium in
HOMOGENEOUS CATALYSIS
the surface layer could be markedly different from the true equilibrium in the gas reaction. It cannot, however, be too strongly emphasized that this does not mean that a catalyst can shift the equilibrium of the gas reaction. The equilibrium ilz the gas phase remains identically the same as it would be if achieved thermally without a catalyst. An analogous case, with two solutions, is that studied by Kuriloff,’ who investigated the equilibrium between @-naphthol and picric acid in water and benzene solutions, in presence of solid picrate. The product of the millimolar concentrations of free naphthol and free undissociated picric acid varied widely in the two solvents, being 2.89 in water and 7550 in benzene, in agreement with the deductions from distribution experiments of the individual substances. The presence of a benzene layer adjacent to the aqueous layer, however, did not in any way disturb the equilibrium in the aqueous layer.
For catalytic reactions in homogeneous systems the inter. mediate compound theory appears to be generally applicable. For most such processes a probable cycle of successive reactions can be postulated. In many cases the intermediate compounds have been isolated. In other cases, the indirect evidence leading to such a conclusion is being steadily brought forward. For example, Jones and Lewis’ give evidence for the formation of an intermediate sucrose-hydrogen-ion complex in the sugar inversion process. In ester hydrolysis the systematic researches of Kendall and his colleagues2 have established the existence of binary and ternary compounds between ester, catalyzing acid, an ‘ water. The tendency towards compound formation is the more marked, the greater the chemical contrast between the basic nature of the ester and the acidity of the catalytic agent. The concordance of this conclusion with the observation that the catalytic activity in ester hydrolysis is greatest with the strong acids and diminkhes with decreasing strength of acid forms a striking piece of evidence in favor of the intermediate compound theory in such systems. Development of the radiation theory of chemical action (Trautz, Lewis, Perrin) has led to the supposition that the necessary energy of reaction is supplied by suitable infra-red radiation. In the beginning, the attempt was made simply to associate the critical energy increment with the heat of reaction and to show that such relationships were plausible in view of the infra-red absorption bands shown by the reacting substances. Recently, Rideal and Hawkinsa have attempted to show that infra-red radiations actually accelerate the velocity of hydrolysis of methyl acetate. A pronounced positive result is claimed. The conclusion, however, can be accepted only with reserve, for the experimental conditions, as far as they may be deduced from the publication, were not ideal. Indeed they were such that, if the positive effect attained is real, the magnitude of the effect of the infra-red radiations must be enormous. The experiments were carried out with IOO cc. of an aqueous solution containing catalyzing acid and ester. The radiation was introduced into the system from above. Owing to the opacity of water to infra-red radiation it is therefore evident that only a film of solution in the surface layer was being irradiated. Since the stirring was only occasional, it is apparent that by far the greater bulk of the solution was not acted upon by the infra-
TABLEIV 1 (Hours)
x(S0s)
ks X 1010
0.5 1.0 1.5 2.0 2.5 3.0
3.5 4.0 5.0
6.0
7 .O 8.0 9.0 10.0 11 .o 12 .o
As a consequence of these considerations it follows that the study of the kinetics of catalytic reactions may give reaction equations totally different from those to be expected from the stoichiometric equation for the gas reaction. This is well known from the experimental work of Fink on the mechanism of the formation of sulfur trioxide from sulfur dioxide and oxygen, of Bodenstein and his co-workers on carbon monoxide and oxygen, and from the recent studies of Armstrong and Hilditch in liquid media. Furthermore, since, as the experiments cited previously show, the distribution of gas between the reaction space and catalyst surface is different a t different partial pressures, it follows that a given equation for the reaction kinetics, while valid over one pressure range, may be invalid over another pressure range. This is clearly shown in many of the kinetic studies
1
2 1 2 . physik.
Chem., 26 (1898), 419.
3
J Chem S o c , 117 (1920). 1120. J . A m Chem S o c , 1914, et seq J . Chem Soc, 117 (1920), 1288.
T H E J O U R N A L O F I N D U S T R I A L A N D E N G I N E E R I N G C H E M I S T R Y . V O ~13, . SO.1
78
red rays. These could be distributed through the solution only by diffusion of the activated hydrogen ions or hydrogen-ionester complexes from the surface into the interior. The question, however, of the possible activity of infra-red rays is so important that duplication and amplification of such
experiments should be undertaken. If such be done the choice of a reaction system through which the radiation might readily penetrate would facilitate the attainment of decisive experimental test. We hope to take such problems in hand a t an early date.
~~
INDUSTRIAL AND AGRICULTURAL CHEMISTRY IN THE. BRITISH WEST INDIES, WITH SOME. ACCOUNT OF THE. WORK OF SIR FRANCIS WATTS, IMPERIAL COMMISSIONER OF AGRICULTURE By C. A. Browne N. Y.SUGARTRADELABORATOKY, 80 SOUTH ST., N E W YORX,N . Y . Received October 5 , 1920
The casual traveler, who makes his first voyage among the West Indian Islands and views from his steamer the crumbling walls of old fortresses, or the remains of stone mansions, acquires at the outset the feeling of a departed civilization. This first impression is intensified by the ruined walls and towers of ancient muscovado sugar works, which, according to the lines of Grainger, the poet of St. Kitts, were once lit up a t night by “far-seen flames bursting through many a chimney.” I t is only when the vessel steams past these scenes of desolation into the harbor of Basseterre, the former home of this poet, and the smoking stacks of a modern sugar factory come into view that the impressions of decadent or vanished industries are dispelled. The present paper is an effort to tell briefly the story of this change from an old t o a new order of things, in which transition CHEMthe efforts of a distinguished member of the AMERICAN ICAL SOCIETY have played a prominent part. With the abolition of slavery in the British West Indies in 1834, the old industrial system of these islands came to an end. The production of sugar, which had always been the chief source of wealth, began to decline, partly from lack of labor and partly from unequal competition with the more scientifically conducted beet-sugar industry of Europe, which marked its phenomenal rise from the date of the abolition of slave labor in the colonies. The inequality of this conflict was later enhanced by the favoring export bounties which beet sugar received, and had it not been for the high prices of sugar, which existed for 2 0 years after the outbreak of the American Civil War, the declining sugar industry of the West Indies would have completely disappeared. The over-stimulation of the beet-sugar industry by bounties and premiums soon had, however, its inevitable effect, and between 1882 and 1892 the price of muscovado fell from 7.3 cents t o 2.8 cents per pound. The industrial condition of the British islands was becoming hopeless, and appeals were made for assistance to the mother country, which for the 50 years followingthe abolition of slavery had shown a strange indifference to its West Indian possessions. This neglect had in fact become so marked that many planters believed their only hope to consist in political union with the United States. It was only with the growing development of the Panama Canal enterprise in the late eighties and the dawning sense of the future strategic and economic importance of the island approaches to this gateway of the Pacific that Great Britain began to take a renewed interest in her tropical colonies. From that time until the present, increasing efforts have been made to improve the industrial, economic, and educational life of the British West Indies. Botanic gardens, experiment stations, and other scientific institutions were established, among the earliest of these being the government laboratory in the island of Antigua, which began its work on Jan. I, 1889, and of which Dr. (now Sir) Francis Watts, a graduate of Mason College, Birmingham, assumed charge as analytical chemist.
IMPROVEMENTS I N SUGAR MANUFACTURE
One of the first investigations which Dr. Watts instituted on beginning his new duties was a thorough examination of the field and factory methods of the sugar industry. His chemical training convinced him t h a t if the cane sugar of the West Indies had to compete with the more scientifically manufactured beet sugar of Europe, the wasteful antiquated processes of the little muscovado factories must disappear. I n a little work, entitled a “Manual for Sugar Growers,” and in various reports, Dr. Watts opened the eyes of the West Indian planters to the enormous losses which their small factory system involved, and as a remedy suggested the erection of large scientifically managed central factories. The idea was favorably received but opinions were divided as to whether such factories should be under government or private control. After much discussion a working scheme was evolved, whereby a group of British capitalists negotiated contracts with certain estate owners in Antigua under which the latter undertook to supply, during a period of 15 years, the sugar canes grown on certain stipulated areas a t a price based on the current market price of sugar, coupled with a share in the profits of the factory and, ultimately, a share in the ownership of the factory itself to the extent of one-half. The capitalists formed a company with a capital of some $200,000, including a sum of $72,000, subscribed by the government. IVith thi? a small central sugar factory
OLD M U S C O V A D O
SUGAR
FACTORY, BRITISHW E S T
INDIES
capable of making about 3000 tons of sugar in a season, was erected a t Gunthorpes, Antigua. The success of the new enterprise was immediate, and the Antigua factory has now grown from a capacity of 3000 to 10,000 tons of sugar per season. In 1919, a t the end of the 15 years’agreement, the government cancelled its $72,000 subscription, its own income from the enterprise in the form of excess profits and exports taxes having exceeded $300,000. The contracting planters received during this time an average of 20 per cent annually on their original investment, and a t the end of the 15 years had turned over to
